Chemical Equilibrium and Le Chatelier's Principle

A state where the forward and reverse reactions occur at the same rate, so the concentrations of reactants and products remain constant.

Describes the relative amounts of reactants and products in an equilibrium mixture.

The position of equilibrium favors the reactants, meaning the equilibrium mixture contains mostly reactants.

A principle stating that when a change is applied to a system at equilibrium, the system will shift to relieve the stress and restore equilibrium.

If temperature is increased, the equilibrium will shift in the endothermic direction to absorb heat and reduce the temperature.

If temperature is decreased, the equilibrium will shift in the exothermic direction to release heat and increase the temperature.

A reaction that can proceed in both the forward and reverse directions.

The process of breaking down a large molecule into smaller units, often with the release of energy.

Adding heat to an exothermic reaction shifts the equilibrium to the left, favoring reactants and reducing product yield. The reaction releases heat to counteract the added heat, trying to restore equilibrium.

Adding heat to an endothermic reaction shifts the equilibrium to the right, favoring products and increasing product yield. The reaction absorbs heat to counteract the added heat, trying to restore equilibrium.

Increasing pressure on a reaction shifts the equilibrium towards the side with fewer moles of gas, reducing the pressure.

Decreasing pressure on a reaction shifts the equilibrium towards the side with more moles of gas, increasing the pressure.

Increasing the concentration of a reactant shifts the equilibrium to the right, favoring product formation. The system tries to consume the added reactant to reach a new balance.

Decreasing the concentration of a reactant shifts the equilibrium to the left, favoring reactant formation. The system attempts to replenish the lost reactant to restore equilibrium.

Changing pressure has no effect on the equilibrium position if the number of moles of gas on both sides of the reaction is the same.

Increasing the concentration of a product shifts the equilibrium to the left, favoring reactant formation. The system attempts to consume the added product to restore equilibrium.

A catalyst speeds up the rate of a reaction by providing an alternate pathway with lower activation energy, but it doesn't change the equilibrium position, because it speeds up both forward and reverse reactions equally.

In industrial processes, the equilibrium position is often adjusted to maximize product yield while considering factors like temperature, pressure, and catalyst use, often involving compromises to balance yield and reaction rate.

The Haber process produces ammonia from nitrogen and hydrogen gas. It is an exothermic reaction, favoring lower temperatures for higher yield, but a compromise temperature is used to balance yield and reaction rate. High pressure favors product formation and is used to increase ammonia yield.

The contact process produces sulfuric acid. It involves two stages, both exothermic. High pressure slightly favors the second stage, increasing yield and rate. A moderate temperature is used to compromise between high yield and reasonable reaction speed.

The hydration of ethene to produce ethanol is an exothermic reaction. Lower temperatures favor the reaction, but a compromise temperature is used to balance yield and reaction rate. Moderate pressure can also be used to slightly favor product formation.

The production of methanol from carbon monoxide and hydrogen gas is an exothermic reaction. Lower temperatures favor the reaction, but a compromise temperature is used to balance yield and reaction rate. Moderate pressure can also be used to increase yield.

The equilibrium constant (Kc) represents the ratio of products to reactants at equilibrium. It is calculated by dividing the product of the equilibrium concentrations of the products raised to their stoichiometric coefficients by the product of the equilibrium concentrations of the reactants raised to their stoichiometric coefficients.

In heterogeneous equilibrium, the concentrations of liquids and solids remain constant and are not included in the Kc expression. This is because their concentrations are independent of the amount present.

The stoichiometric coefficients (m, n, p, q) in a balanced chemical equation represent the number of moles of each species involved in the reaction. They are used to determine the exponents in the Kc expression.

The reaction rate is the speed at which a reaction proceeds. It is influenced by factors like temperature, concentration, and the use of catalysts.

A catalyst is a substance that speeds up a reaction without being consumed in the process. Catalysts lower the activation energy required for a reaction to occur.

Higher pressure generally favors a reaction that produces fewer moles of gas. This is because the system seeks to minimize the pressure by shifting towards the side with fewer gas molecules.

Temperature affects the equilibrium constant (Kc). Increasing temperature favors an endothermic reaction (heat absorbed), while decreasing temperature favors an exothermic reaction (heat released).

Recycling unreacted reactants back into the reactor can improve the overall yield of a reaction by maximizing the conversion of reactants to products.